CHEM 1305 Final Exam Review: Comprehensive Study Notes for General Chemistry

Electron Configuration

Understanding Electron Arrangement in Atoms

• Electron configuration describes the distribution of electrons among the atomic orbitals of an element.

• Electrons fill orbitals in order of increasing energy, following the Aufbau principle, Pauli exclusion principle, and Hund's rule.

• Example: Silicon (Si, atomic number 14) has the configuration $1s^2 2s^2 2p^6 3s^2 3p^2$.

Isotopes & Ions

Atomic Structure and Charge

• Isotopes are atoms of the same element with different numbers of neutrons.

• Ions are atoms or molecules with a net electric charge due to loss or gain of electrons.

• Example: An atom with 20 protons, 18 electrons, and 20 neutrons is a $Ca^{2+}$ ion.

Valence Electrons

Outer Shell Electrons and Chemical Reactivity

• Valence electrons are electrons in the outermost shell, crucial for bonding and chemical reactions.

• Example: Oxygen (O) has 6 valence electrons.

Bond Type

Classification of Chemical Bonds

• Ionic bonds form between metals and nonmetals via electron transfer.

• Covalent bonds involve electron sharing between nonmetals.

• Metallic bonds occur between metals, with delocalized electrons.

• Example: Cl–Cl is a nonpolar covalent bond.

Oxidation State

Assigning Oxidation Numbers

• Oxidation state indicates the hypothetical charge an atom would have if all bonds were ionic.

• Example: In $Fe_2(SO_4)_3$, iron has an oxidation state of +3.

Formula Writing

Constructing Chemical Formulas

• Formulas represent the types and numbers of atoms in a compound.

• Example: Calcium oxide is $CaO$.

Oxidation-Reduction (Redox)

Electron Transfer Reactions

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Example: In $Zn + Cu^{2+} \rightarrow Zn^{2+} + Cu$, Zn is oxidized, Cu2+ is reduced.

Reaction Type

Classifying Chemical Reactions

• Combination: Two or more substances form one product.

• Decomposition: One substance breaks into two or more products.

• Single Replacement: One element replaces another in a compound.

• Double Replacement: Exchange of ions between two compounds.

• Example: $Fe + CuSO_4 \rightarrow FeSO_4 + Cu$ is single replacement.

Balancing

Conservation of Mass in Chemical Equations

• Balancing ensures equal numbers of each atom on both sides of the equation.

• Example: $N_2 + 3H_2 \rightarrow 2NH_3$

Stoichiometry

Quantitative Relationships in Reactions

• Stoichiometry uses balanced equations to calculate amounts of reactants and products.

• Example: $2NH_3 \rightarrow 2N_2 + 3H_2$

• Use molar mass and mole ratios for calculations.

STP Moles

Standard Temperature and Pressure Calculations

• At STP (0°C, 1 atm), 1 mole of gas occupies 22.4 L.

• Example: 22.4 L of gas at STP contains 1 mole.

Gas Mole Fraction

Relative Amounts of Gases in Mixtures

• Mole fraction ($X$) = moles of component / total moles.

• Example: If 2 white and 8 gray gas particles, $X_{white} = 2/10 = 0.2$.

Gas Law

Relationships Between Pressure, Volume, and Temperature

• Combined Gas Law: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

• Example: Calculate new pressure or volume when temperature changes.

Kinetic Molecular Theory (KMT)

Model for Gas Behavior

• Gases consist of particles in constant, random motion.

• Collisions are elastic; no energy is lost.

• No intermolecular forces; volume of particles negligible.

Solute/Solvent

Components of Solutions

• Solute: Substance dissolved.

• Solvent: Substance doing the dissolving (usually present in greater amount).

• Example: In saltwater, salt is solute, water is solvent.

Electrolytes

Conductivity in Solutions

• Electrolytes dissociate into ions in solution, conducting electricity.

• Strong electrolytes: Complete dissociation (e.g., NaCl).

• Weak electrolytes: Partial dissociation (e.g., acetic acid).

• Nonelectrolytes: Do not dissociate (e.g., sugar).

Dilution

Preparing Solutions of Lower Concentration

• Use $M_1V_1 = M_2V_2$ to calculate new concentration after dilution.

• Example: Diluting 25.0 mL of 4.0 M NaOH to 100.0 mL gives $M_2 = 1.0$ M.

Net Ionic Equation

Representing Only the Species Involved in Reaction

• Net ionic equations show only ions and molecules directly involved in the reaction.

• Example: $Ag^+ (aq) + Cl^- (aq) \rightarrow AgCl (s)$

Thermochemistry

Heat Changes in Chemical Reactions

• Endothermic: Absorbs heat ($\Delta H > 0$).

• Exothermic: Releases heat ($\Delta H < 0$).

• Example: Dissolving NH4NO3 in water is endothermic.

IUPAC Naming

Systematic Naming of Compounds

• IUPAC rules provide standardized names for chemical compounds.

• Example: $H_2O$ is dihydrogen monoxide.

ΔH Reaction

Enthalpy Change in Reactions

• $\Delta H$ is the heat change at constant pressure.

• Example: $H_2O(l) \rightarrow H_2O(g)$ is endothermic ($\Delta H > 0$).

Heat Equation

Calculating Heat Transfer

• Use $q = mC\Delta T$ where $q$ is heat, $m$ is mass, $C$ is specific heat, $\Delta T$ is temperature change.

• Example: To heat 540 g from 25°C to 75°C, $q = 540 \times 0.44 \times (75-25)$

Hydrogen Composition

Hydrogen Content in Compounds

• Compare number of hydrogen atoms in different molecules.

• Example: $C_2H_6$ has more hydrogen than $C_2H_2$.

Periodic Property

Trends in the Periodic Table

• Properties such as atomic radius, ionization energy, electron affinity, and electronegativity show periodic trends.

• Example: Atomic radius decreases across a period, increases down a group.

Molecular Geometry

Shapes of Molecules

• VSEPR theory predicts molecular shapes based on electron pair repulsion.

• Examples: Tetrahedral ($CH_4$), Trigonal planar ($BF_3$), Bent ($H_2O$).

Formal Charge

Determining Charge Distribution in Molecules

• Formal charge = (valence electrons) - (nonbonding electrons) - ½(bonding electrons).

• Example: In $NH_4^+$, N has a formal charge of +1.

Electrostatic Map

Visualizing Charge Distribution

• Electrostatic maps show regions of partial positive and negative charge in molecules.

• Example: In $H_2O$, O is partially negative, H is partially positive.

Photoelectric Effect

Emission of Electrons from Metals

• Electrons are emitted when light of sufficient frequency strikes a metal surface.

• Increasing frequency (shorter wavelength) increases electron emission.

Light Frequency

Calculating Frequency of Electromagnetic Radiation

• Frequency ($\nu$) is related to wavelength ($\lambda$) by $c = \lambda \nu$ where $c$ is speed of light.

• Example: Orange light ($\lambda = 620$ nm) has $\nu = c/\lambda$.

Bohr Model

Energy Levels in Hydrogen Atom

• Energy of electron in hydrogen: $E_n = -2.18 \times 10^{-18} \frac{1}{n^2}$ J

• Example: Given $E = -3.40 \times 10^{-19}$ J, solve for $n$.

Topic	Key Formula/Concept	Example
Gas Law	$\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$	Calculate new pressure after temperature change
Heat Equation	$q = mC\Delta T$	Find heat needed to warm a substance
STP Moles	1 mole gas = 22.4 L at STP	Convert volume to moles
Formal Charge	valence - nonbonding - ½(bonding)	Calculate for $NH_4^+$

Additional info: Some explanations and examples have been expanded for clarity and completeness beyond the original review questions.